collision theory

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Reaction Rates Prepared by Barbara Andreswith A. NiedzvieckiApril 5, 2004@ Centennial Regional High School : 

Reaction Rates Prepared by Barbara Andreswith A. NiedzvieckiApril 5, 2004@ Centennial Regional High School

Factors which influence reaction rates: : 

Factors which influence reaction rates: Temperature Presence of a catalyst Concentration of reactants Surface area of reactants Type of reactants

How can we explain these factors? : 

How can we explain these factors?

Collision Theory : 

Collision Theory

Reaction rate depends on: : 

Reaction rate depends on: # of collision per unit time Success rate

Success depends on the: : 

Success depends on the: Collision geometry (orientation) Energy of the collision

Collision Geometry : 

Collision Geometry

H2(g) + I2(g) 2HI(g) Side-to-Side Collision : 

H2(g) + I2(g) 2HI(g) Side-to-Side Collision

Slide 9: 

H2(g) + I2(g) X 2HI(g) End-to-End Collision

Slide 10: 

H2(g) + I2(g) X 2HI(g) Glancing Collision

HCl(g) + C2H4(g) products : 

HCl(g) + C2H4(g) products For this example: Hydrogen end of H-Cl approaches C=C double bond ….. Success All other collisions …..“rebound”

Which of the following collisions will be successful? : 

Which of the following collisions will be successful?  X X X

Energy of the Collision : 

Energy of the Collision

Slide 14: 

Particles (molecules & atoms) must collide with a minimum energy Minimum energy called the activation energy of the reaction

Remember, the energy profile for an exothermic reaction looks like this: : 

Remember, the energy profile for an exothermic reaction looks like this:

In order for a reaction to occur the: : 

In order for a reaction to occur the: activation energy “barrier” must be overcome collision energy  activation energy bond breaking & bond formation (reaction) collision energy < activation energy “rebound” (no reaction)

For any reacting system: : 

For any reacting system: Particles have a very wide range of energies What proportion of the particles have high enough energies to react when they collide?

Slide 18: 

Maxwell-Boltzmann Distribution (gases) Area under the curve = total number of particles present

Slide 19: 

Maxwell-Boltzmann Distribution (gases)

To enable low energy particles to react we have to: : 

To enable low energy particles to react we have to: change the shape of the curve or move the activation energy to the left

How can we change the shape of the curve? : 

How can we change the shape of the curve? Change the temperature of the reaction

Slide 22: 

How can we overcome the activation energy “barrier”?

Slide 23: 

Effect of Temperature on Reaction Rates

Temperature  Reaction Rate  : 

Temperature  Reaction Rate  For many reactions happening at around room temperature, the rate of reaction doubles for every 10°C rise in temperature.

Why does this occur? : 

Why does this occur? Increased collision frequency Particles only react when they collide. Heat a substance … the particles move faster Collide more frequently Increased rate of reaction

Slide 26: 

Increased kinetic energy Overcome activation energy “barrier” Heat a substance … the particles move faster Collisions are more “intense” Increased rate of reaction.

Temperature  … graph shape altered : 

Temperature  … graph shape altered Area under a curve = count of # of particles “T + t” area doubled Doubled rate of reaction

Slide 28: 

Effect of a Catalyst on Reaction Rates

Catalysts are substances that: : 

Catalysts are substances that: speed up a reaction participate in the reaction steps are chemically unchanged at the end of the reaction have no loss of mass after the reaction

Catalysts … : 

Catalysts … provide an alternative way for the reaction to happen which has a lower activation energy move the activation energy to the left on Maxwell-Boltzmann Distribution graph.

Maxwell-Boltzmann Distribution : 

Maxwell-Boltzmann Distribution

On an energy profile: : 

On an energy profile:

Slide 33: 

Effect of Concentration on Reaction Rates

Slide 34: 

Na2S2O3(aq) + 2 HCl  S(s) + SO2(aq) + 2 NaCl(aq) + H2O(l) Dilute HCl is added to sodium thiosulphate solution …. precipitate of sulphur forms Sodium thiosulphate solution diluted …. precipitate takes longer and longer to form.

Case 1: Reactions - 2 Particles : 

Case 1: Reactions - 2 Particles Must first collide Concentration is higher  the chance of collision are greater.

Slide 36: 

Case 2: Catalyst already working as fast as it can Small amount of a solid catalyst in a reaction High concentration of reactants Catalyst completely “occupied” Adding more reactants has no effect

Slide 37: 

Case 3: Certain Multi-step Reactions

Slide 38: 

Pressure  Concentration  Reaction Rate  Mass of gas … squeezed into a smaller volume Relationship Between Pressure and Concentration

Can also be explained by the ideal gas law … : 

Can also be explained by the ideal gas law …

Slide 40: 

Effect of Surface Area on Reaction Rates

Surface Area  Reaction Rate  : 

Surface Area  Reaction Rate  Particles in the gas or liquid must collide with the solid particles Increasing the surface area of the solid increases the chances of collision taking place

Slide 42: 

2Mg(s)+ 2 HCl(aq)  H2(g)+2 MgCl

Slide 43: 

Effect of Nature of Reactants on Reaction Rates

We have studied the following two reactions: : 

We have studied the following two reactions: C25H52(s) + 38 O2(g)  25 CO2(g) + 26 H2O(g)  2 CH3OH(l) + 3 O2(g)  2 CO2(g) + 4 H2O(g) Based on experimental results, how did the nature of reactants affects the rate of reaction: # of bond broken and formed  reaction rate 

We have studied the following two reactions: : 

We have studied the following two reactions: C25H52(s) + 38 O2(g)  25 CO2(g) + 26 H2O(g)  2 CH3OH(l) + 3 O2(g)  2 CO2(g) + 4 H2O(g) Explain this based on Collision Theory

We have studied the following two reactions: : 

We have studied the following two reactions: C25H52(s) + 38 O2(g)  25 CO2(g) + 26 H2O(g)  2 CH3OH(l) + 3 O2(g)  2 CO2(g) + 4 H2O(g) Explain this based on Collision Theory Greater # of collisions required

The End : 

The End

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